Video transcript

Electronegativity is probably
the most important concept to understand
inorganic chemistry. We're going to use a definition
that Linus Pauling gives in his book, The Nature
of the Chemical Bond. So Linus Pauling says
that electronegativity refers to the power of
an atom in a molecule to attract electrons to itself. So if I look at a
molecule, I'm going to compare two atoms
in that molecule. I'm going to compare
carbon to oxygen in terms of the
electronegativity. And to do that, I need to
look over here in the right at the organic periodic table,
which shows the elements most commonly used in
organic chemistry. And then in blue, it
gives us the Pauling scale for electronegativity. So Linus Pauling
actually calculated electronegativity
values for the elements and put them into the table. And that allows us to
compare different elements in terms of their
electronegativities. For example, we are
concerned with carbon, which has an electronegativity
value of 2.5. And we're going to
compare that to oxygen, which has an electronegativity
value of 3.5. So oxygen is more
electronegative than carbon. And the definition
tells us that if oxygen is more electronegative,
oxygen has a greater power to attract electrons to
itself than carbon does. And so if you think
about the electrons and the covalent bonds
between carbon and oxygen that are shared, they're
shared unequally. Because oxygen is
more electronegative, oxygen is going to
pull those electrons in red closer to itself. And since electrons
are negatively charged, the oxygen is going
to get a little bit more negative charge. And so it's going
to have what we call a partial
negative charge on it. So partial negative. Its partial sign is a
lowercase Greek letter, delta. And so the oxygen is
partially negative. It's pulling the electrons
in red closer to itself. Another way to show the movement
of those electrons in red closer to the oxygen would
be this funny arrow here. So the arrow points
in the direction of the movement of
the electrons in red. So carbon is losing some
of those electrons in red. Carbon is losing a little
bit of electron density. Carbon is losing a little
bit of negative charge. So carbon used to be
neutral, but since it's losing a little bit
of negative charge, this carbon will end up being
partially positive, like that. So the carbon is
partially positive. And the oxygen is
partially negative. That's a polarized situation. You have a little bit
of negative charge on one side, a little
bit of positive charge on the other side. So let's say it's
still a covalent bond, but it's a polarized
covalent bond due to the differences
in electronegativities between those two atoms. Let's do a few
more examples here where we show the differences
in electronegativity. So if I were thinking
about a molecule that has two carbons
in it, and I'm thinking about what happens
to the electrons in red. Well, for this
example, each carbon has the same value
for electronegativity. So the carbon on the
left has a value of 2.5. The carbon on the right
has a value of 2.5. That's a difference in
electronegativity of zero. Which means that
the electrons in red aren't going to move
towards one carbon or towards the other carbon. They're going to
stay in the middle. They're going to be shared
between those two atoms. So this is a covalent bond, and
there's no polarity situation created here since
there's no difference in electronegativity. So we call this a
non-polar covalent bond. This is a non-polar
covalent bond, like that. Let's do another example. Let's compare
carbon to hydrogen. So if I had a
molecule and I have a bond between
carbon and hydrogen, and I want to know what
happens to the electrons in red between the carbon and hydrogen. We've seen that. Carbon has an
electronegativity value of 2.5. And we go up here to hydrogen,
which has a value of 2.1. So that's a difference of 0.4. So there is the difference
in electronegativity between those two atoms, but
it's a very small difference. And so most textbooks
would consider the bond between
carbon and hydrogen to still be a non-polar
covalent bond. All right. Let's go ahead and
put in the example we did above, where we compared
the electronegativities of carbon and oxygen, like that. When we looked up
the values, we saw that carbon had an
electronegativity value of 2.5 and oxygen had a value of
3.5, for difference of 1. And that's enough to have
a polar covalent bond. Right? This is a polar covalent
bond between the carbon and the oxygen. So when we think about
the electrons in red, electrons in red are pulled
closer to the oxygen, giving the oxygen a
partial negative charge. And since electron density is
moving away from the carbon, the carbon gets a
partial positive charge. And so we can see that if your
difference in electronegativity is 1, it's considered to
be a polar covalent bond. And if your difference in
electronegativity is 0.4, that's considered to be a
non-polar covalent bond. So somewhere in
between there must be the difference between
non-polar covalent bond and a polar covalent bond. And most textbooks will tell
you approximately somewhere in the 0.5 range. So if the difference
in electronegativity is greater than 0.5,
you can go ahead and consider it to be mostly
a polar covalent bond. If the difference
in electronegativity is less than 0.5,
we would consider that to be a non-polar
covalent bond. Now, I should point
out that we're using the Pauling scale
for electronegativity here. And there are several different
scales for electronegativity. So these numbers
are not absolute. These are more
relative differences. And it's the relative
difference in electronegativity that we care the most about. Let's do another example. Let's compare
oxygen to hydrogen. So let's think
about what happens to the electrons between
oxygen and hydrogen. So the electrons in red here. All right. So we've already seen the
electronegativity values for both of these atoms. Oxygen had a value of 3.5, and
hydrogen had a value of 2.1. So that's an electronegativity
difference of 1.4. So this is a polar
covalent bond. Since oxygen is more
electronegative than hydrogen, the electrons in red are going
to move closer to the oxygen. So the oxygen is going to get
a partial negative charge. And the hydrogen is going to
get a partial positive charge, like that. All right. Let's do carbon and lithium now. So if I go ahead and draw a
bond between carbon and lithium, and once again, we are
concerned with the two electrons between carbon and lithium. The electronegativity value
for carbon we've seen is 2.5. We need to go back up
to our periodic table to find the electronegativity
value for lithium. So I go up here,
and I find lithium in group one of
my periodic table has an electronegativity
value of 1. So I go back down here, and
I go ahead and put in a 1. And so that's a difference
in electronegativity of 1.5. So we could consider this
to be a polar covalent bond. This time, carbon is more
electronegative than lithium. So the electrons
in red are going to move closer to
the carbon atom. And so the carbon is going to
have a little bit more electron density than usual. So it's going to be
partially negative. And the lithium is
losing electron density, so we're going to say that
lithium is partially positive. Now here, I'm treating this
bond as a polar covalent bond. But you'll see in a few
minutes that we could also consider this to
be an ionic bond. And that just depends on
what electronegativity values you're dealing with,
what type of chemical reaction that you're working with. So we could consider
this to be an ionic bond. Let's go ahead and do
an example of a compound that we know for sure is ionic. Sodium chloride, of course,
would be the famous example. So to start with,
I'm going to pretend like there's a covalent
bond between the sodium and the chlorine. All right. So I'm going to say there's a
covalent bond to start with. And we'll put in our electrons. And we know that this bond
consists of two electrons, like that. Let's look at the differences
in electronegativity between sodium and chlorine. All right. So I'm going to go back up here. I'm going to find sodium,
which has a value of 0.9, and chlorine which
has a value of 3. So 0.9 for sodium
and 3 for chlorine. So sodium's value is 0.9. Chlorine's is 3. That's a large difference
in electronegativity. That's a difference of 2.1. And so chlorine is much more
electronegative than sodium. And it turns out, it's so
much more electronegative that it's no longer going to
share electrons with sodium. It's going to steal
those electrons. So when I redraw
it here, I'm going to show chlorine being
surrounded by eight electrons. So these two electrons in red--
let me go ahead and show them-- these two electrons in red
here between the sodium and the chlorine, since
chlorine is so much more electronegative, it's going
to attract those two electrons in red so strongly that
it completely steals them. So those two
electrons in red are going to be stolen by
the chlorine, like that. And so the sodium
is left over here. And so chlorine has
an extra electron, which gives it a
negative 1 formal charge. So we're no longer talking
about partial charges here. Chlorine gets a full
negative 1 formal charge. Sodium lost an
electron, so it ends up with a positive formal
charge, like that. And so we know this is an ionic
bond between these two ions. So this represents
an ionic bond. So the difference
in electronegativity is somewhere
between 1.5 and 2.1, between a polar covalent
bond and an ionic bond. So most textbooks we'll
see approximately somewhere around 1.7. So if you're higher
than 1.7, it's generally considered to
be mostly an ionic bond. Lower than 1.7, in the
polar covalent range. But that doesn't always
have to be the case. Right? So we'll come back
now to the example between carbon and lithium. So if we go back up here
to carbon and lithium, here we treat it like
a polar covalent bond. But sometimes you might want
to treat the bond in red as being an ionic bond. So let's go ahead and draw a
picture of carbon and lithium where we're treating
it as an ionic bond. So if carbon is more
electronegative than lithium, carbon's going to steal
the two electrons in red. So I'll go ahead and
show the electrons in red have now moved on
to the carbon atom. So it's no longer sharing
it with the lithium. Carbon has stolen
those electrons. And lithium is over here. So lithium lost one of its
electrons, giving it a plus 1 formal charge. Carbon gained an
electron, giving it a negative 1 formal charge. And so here, we're treating
it like an ionic bond. Full formal charges here. And this is useful for some
organic chemistry reactions. And so what I'm trying
to point out here is these divisions,
1.7, it's not absolute. It's a relative thing. You could draw the
dot structure above, and this would be
considered be correct. Right? You could draw it like this. Or you could treat it like
an ionic bond down here. This is relatively
close to the cutoff. So this is an overview
of electronegativity. And even though
we've been dealing with numbers in this
video, in future videos, we don't care so much
about the numbers. We care about the
relative differences in electronegativity. So it's important to understand
something as simple as oxygen is more electronegative
than carbon. And that's going to
help you when you're doing organic
chemistry mechanisms.